Chapter 8 Notes: Basic Concepts of Chemical Bonding

Chemical Bonds

• There are three basic types of bonds:
  • Ionic: Electrostatic attraction between ions.
  • Covalent: Sharing of electrons.
  • Metallic: Free electrons hold metal atoms together.

Lewis Symbols and the Octet Rule

• G. N. Lewis developed a method to denote potential bonding electrons by using one dot for every valence electron around the element symbol.
• When forming compounds, atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons (the octet rule).

Ionic Bonding

• Occurs between metals and nonmetals (except group 8A).
• Involves electron transfer.
• Very exothermic.
• One element readily gives up an electron (has a low ionization energy).
• Another element readily gains an electron (has a high electron affinity).
• Arrow(s) indicate the transfer of the electron(s).

Properties of Ionic Substances

• Evidence of a well-defined three-dimensional structure:
  • Brittle
  • High melting points
  • Crystalline
  • Cleave along smooth lines

Energetics of Ionic Bonding

• Many factors affect the energy of ionic bonding.
• Start with the metal and nonmetal elements:
  • Make gaseous atoms.
  • Make ions.
  • Combine the ions: $NaCl(s)$. The Born–Haber Cycle
• Chapter 7 presented ion formation (ionization energy and electron affinity).
• Energy is required to convert the elements to atoms (endothermic).
• Energy is required to create a cation (endothermic).
• Energy is released to form the anion (exothermic).
• Formation of a solid releases a huge amount of energy (exothermic).
• This makes the formation of salts from the elements exothermic.

Lattice Energy

• This is a measure of how much stability results in arranging oppositely charged ions in an ionic solid.
• Defined as the energy required to completely separate one mole of a solid ionic compound into its gaseous ions.
• Determined using the Born-Haber Cycle calculation.
• This is designated as the amount of energy released to form the ionic compound (Table 8.1).
• Lattice energy increases with:
  • Increasing charge $(Q)$ on the cations.
  • Decreasing size $(d)$ of the ions.

Electron Configuration of Ions

• Main group metals lose electrons, resulting in the electron configuration of the previous noble gas.
• Nonmetals gain electrons, resulting in the electron configuration of the nearest noble gas.
• Transition metals do not follow the Octet rule.
• Transition metals lose the valence electrons first (highest $n$), then lose the $d$-electrons necessary for the given ion charge.

Covalent Bonding

• In covalent bonds, atoms share electrons (primarily nonmetals).
• Electrostatic interactions in these bonds:
  • Attractions between electrons and nuclei
  • Repulsions between electrons
  • Repulsions between nuclei
• For a bond to form, the attractions must be greater than the repulsions.

Lewis Structures

• Sharing electrons to make covalent bonds can be demonstrated using Lewis structures.
• We start by trying to give each atom the same number of electrons as the nearest noble gas by sharing electrons.
• The simplest examples are for hydrogen, and chlorine, shown below.

Number of Bonds for Nonmetals

• The group number is the number of valence electrons.
• To get an octet (nearest noble gas configuration) in the simplest covalent molecules for nonmetals, the number of bonds needed will be the same as the electrons needed to complete the octet.

Electrons Using Lewis Structures

• Lone pairs: an unshared pair of electrons located on only one atom in a Lewis structure.
• Bonding pairs: shared electrons in a Lewis structure between two atoms. Can be represented either by two dots or one line.

Multiple Bonds

• Single bonds: atoms share only one pair of electrons.
• Double bonds: two pairs of electrons are shared.
• Triple bonds: three bonds are shared between two atoms.

Polarity of Bonds and Electronegativity

• The electrons in a covalent bond are not always shared equally.
• Bond polarity is a measure of how equally or unequally the electrons in a covalent bond are shared.
• In a nonpolar covalent bond, the electrons are shared equally.
• In a polar covalent bond, one of the atoms attracts electrons to itself with a greater force.

Polar or Nonpolar Covalent Bonds

• In elemental fluorine, the atoms pull electrons equally. The bond is a nonpolar covalent bond.
• In hydrogen fluoride, fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does. Therefore, the fluorine end of the molecule has more electron density than the hydrogen end, making it a polar covalent bond. LiF has ionic bonding.

Electronegativity

• Electronegativity is the ability of an atom in a molecule to attract electrons to itself.
• On the periodic table, electronegativity generally increases as you go:
  • From left to right across a period.
  • From the bottom to the top of a group.

Electronegativity and Polar Covalent Bonds

• When two atoms share electrons unequally, a polar covalent bond results.
• Electrons tend to spend more time around the more electronegative atom. The result is a partial negative charge (not a complete transfer of charge) represented by $\delta$-.
• The other atom is “more positive”.

Polar Covalent Bonds

• The greater the difference in electronegativity, the more polar is the bond.

Dipoles

• When two electrical charges of equal magnitude, but opposite sign, are separated by a distance, a dipole forms.
• A dipole moment, $\mu$, produced by two equal but opposite charges separated by a distance, $r$, is calculated: $\mu = Qr$
• It is measured in debyes (D).

Comparing Ionic and Covalent Bonding

• There is a gradual transition between the two types of bonding.
  • Ionic: complete electron transfer (metal + nonmetal).
  • Covalent: electron pair sharing (two nonmetals).
• Use electronegativity difference to guide:
  • Electronegative difference > 2.0 is often ionic
  • Exceptions exist. Take into account the metal oxidation number. Higher oxidation numbers can be covalent.
• Also look at physical properties of compounds (lower melting points mean covalent bonding).

Drawing Lewis Structures

• Sum the valence electrons from all atoms, taking into account overall charge.
  • If it is an anion, add one electron for each negative charge.
  • If it is a cation, subtract one electron for each positive charge.

Writing Lewis Structures

• Write the symbols for the atoms, show which atoms are attached to which, and connect them with a single bond (a line representing two electrons).
• Complete the octets around all atoms bonded to the central atom.
• Place any remaining electrons on the central atom.
• If there are not enough electrons to give the central atom an octet, try multiple bonds.
• Then assign formal charges.
• Formal charge is the charge an atom would have if all of the electrons in a covalent bond were shared equally.
• The dominant Lewis structure:
  • is the one in which atoms have formal charges closest to zero.
  • puts a negative formal charge on the most electronegative atom.
• As such, it can be used to decide which structure is best.

The Best Lewis Structure?

• Following the rules, there are two Lewis structures for ozone.
• However, it doesn’t agree with what is observed in nature: Both O-to-O bond lengths are the same.

Resonance Structures

• We use multiple structures, resonance structures, to describe the molecule.
• It is used when one Lewis structure cannot accurately depict a molecule like ozone.
• For ozone, the two bond lengths are equivalent and intermediate between a single and double bond.
• The organic compound benzene, has two resonance structures. Experiment shows six equivalent bond lengths.
• Benzene is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.
• Localized electrons are specifically on one atom or shared between two atoms. Delocalized electrons are shared by multiple atoms.

Exceptions to the Octet Rule

• There are three types of ions or molecules that do not follow the octet rule:
  • Ions or molecules with an odd number of electrons
  • Ions or molecules with less than an octet
  • Ions or molecules with more than eight valence electrons (an expanded octet)

Odd Number of Valence Electrons

• Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons, i.e., nitrogen oxide.

Fewer Than Eight Valence Electrons

• Elements in the second period before carbon can make stable compounds with fewer than eight electrons.
• Consider $\text{BF}_3$:
  • Giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine.
  • This would not be an accurate picture of the distribution of electrons in $\text{BF}_3$.
• If filling the octet of the central atom results in a negative formal charge on the central atom and a positive formal charge on the more electronegative outer atom, don’t fill the octet of the central atom.

More Than Eight Valence Electrons

• When an element is in periods 3 through 6, it can use d- orbitals to make more than four bonds and be hypervalent.
• Examples: $\text{SF}_6$ and the phosphate ion (Note: Phosphate will actually have four resonance structures with five bonds on the P atom.)

Strengths and Lengths of Single and Multiple Bonds

• Introduced for single covalent bonds in Chapter 5. Extend now to multiple covalent bonds. All bond enthalpies are positive as bond breaking is an endothermic process.
• These are averages over many compounds; not every bond for a pair of atoms has the same bond energy.

Bond Enthalpy and Bond Length

• Multiple bonds are stronger than single bonds.
• As the number of bonds between two atoms increases, the bond length decreases.